Thermodynamics electromotive force

The concentration overpotential i/c is the component of the overpotential due to concentration gradients in the electrolyte solution near the electrode, not including the electric double layer. The concentration overpotential is usually identified with the Nernst potential of the working electrode with respect to the reference electrode that is, the thermodynamic electromotive force (emf) of a concentration cell formed between the working electrode (immersed in electrolyte depleted of reacting species) and the reference electrode (of the same kind but immersed in bulk electrolyte solution) ... 

The different fuel-cell systems differ in the nature of the components selected, and thus in the nature of the current-producing chemical reaction. Each reaction is associated with a particular value of enthalpy and Gibbs free energy -AG) of the reaction and thus also with a particular value of the heat of reaction and of the thermodynamic electromotive force (EMF) e. 

The thermodynamic electromotive force of a polymer electrolyte membrane fuel cell at a temperature of 25°C is given by e = 1.229 V. The open-circuit voltage (OCV) of a hydrogen-oxygen polymer electrolyte membrane fuel cell has values between 0.95 and 1.02 V, depending on the temperature and gas pressures. 

The thermodynamic electromotive force of MCFC, called the open circuit voltage ( ocv)> is determined by the following equation according to Eqs. 8.1a and 8.1b ... 

Values taken from S. Glasstone. Thermodynamics for Chemists. D. Van Nostrand Company Inc., Toronto, p. 443 (1947). The values tabulated in this reference were taken from D. N. Craig and G. W. Vinal, J. Res. Natl. Bur. Stand.. Thermodynamic Properties of Sulfuric Acid Solutions and Their Relation to the Electromotive Force and Heat of Reaction of the Lead Storage Battery", 24, 475-490 (1940). More recent values at the higher molality can be found in W. F. Giauque. E. W. Hornung. J. E. Kunzler and T. R. Rubin, The Thermodynamic Properties of Aqueous Sulfuric Acid Solutions and Hydrates from 15 to 300° K", J. Am. Chem. Soc.. 82, 62-70 (1960). 

To recall the basic concepts of the thermodynamics of cell operation, such as the electrode potential E, the standard electrode potential and the electromotive force (emf). 

The portion AQ = AH - AG = TAS of AH is transformed into heat. Ideal theoretical efficiencies % determined by the types and amounts of reactants and by the operating temperature. Fuel cells have an efficiency advantage over combustion engines because the latter are subdued to the Carnot limitation. High thermodynamic efficiencies are possible for typical fuel cell reactions (e.g., e,h = 0.83 (at 25°C) for H2 + I/2O2 -> H20(i)). The electrical potential difference between anode and cathode, = -AG/W(f, which is also called the electromotive force or open-circuit voltage, drives electrons through the external... 

The thermodynamic equation that expresses the electromotive force E in terms of the standard potential E° and the concentrations of reactants and products in a redox reaction. 

The electromotive force (EMF) generated by electrochemical cells can be used to measure partial Gibbs energies which, like vapour pressure measurements, distinguishes these methods from other techniques that measure integral thermodynamic quantities. Following Moser (1979), a typical cell used to obtain results on Zn-ln-Pb is represented in the following way ... 

As we have seen, acidity and basicity are intimately connected with electron transfer. When the electron transfer involves an integral number of electrons it is customary to refer to the process as a redox reaction. This is not the place for a thorough discussion of the thermodynamics of electrochemistry that may be found in any good textbook of physical chemistry. Rather, we shall investigate the applications of electromotive force (emf) of interest to the inorganic chemist. Nevertheless, a very brief review of the conventions and thermodynamics of electrode potentials and half-reactions will be presented. 

With an understanding of the meaning and measurement of the difference of electrical potential, we can develop the thermodynamics of a galvanic cell. We choose a specific cell, but one in which many of the principles related to the obtaining of thermodynamic data from measurement of the electromotive forces (emf) of the cell are illustrated. The specific cell is depicted as... 

Obviously, plasmas can be used very efficiently within the synthetic approach (i), and all examples given in this paper are assigned to the synthetic approach. It is much less obvious whether plasmas can be used also in the counter-direction. In order to measure a stable and reproducible electromotive force (EMF) the corresponding electrochemical (galvanic) cell must be in (local) thermodynamic equilibrium. Low-temperature plasmas represent non-equilibrium states and are highly inhomogeneous systems from a thermodynamic point of view, often not... 

Guggenheim has expressed1 the electromotive force of the cell (C. 1) in terms of the thermodynamic electrochemical potentials . The final result... 

A more complete description of this research will soon appear in the Journal of the American Chemical Society. The preparation of the cells so as to secure constancy and reprodudbility of the electromotive force values, the methods of making the measurements, the full experimental data, and thermodynamic calculations from them of other free-energy values will be there presented in detail. 

Figure 7.5 shows a composed of a bimetallic couple metal wires a and b with one junction maintained at temperature T and the other maintained at T+ dT. An electromotive force E causes a current / to pass through the wires. A Peltier heat qpe(T + dT) per unit current will be absorbed at the warm junction and an amount of heat qpe(T) will be given off at the cool junction. To maintain a temperature gradient, Thomson heat (q l h i)(dT) must be supplied to the metal a, and an amount of heat (r/Th h)(c/7 j must be removed from b, since the current is in the opposite direction in metal wire b. In a closed work cycle, the electric energy is fully converted to heat. Therefore, the energy balance per unit current by the first law of thermodynamics is... 

Electromotive Forces and Thermodynamic Functions of the Cell Pt, H2 HBr(m), X% Alcohol, Y% Water AgBr-Ag in Pure and Mixed Solvents... 

Electromotive force measurements of the cell Pt, H2 HBr(m), X% alcohol, Y% water AgBr-Ag were made at 25°, 35°, and 45°C in the following solvent systems (1) water, (2) water-ethanol (30%, 60%, 90%, 99% ethanol), (3) anhydrous ethanol, (4) water-tert-butanol (30%, 60%, 91% and 99% tert-butanol), and (5) anhydrous tert-butanol. Calculations of standard cell potential were made using the Debye-Huckel theory as extended by Gronwall, LaMer, and Sandved. Gibbs free energy, enthalpy, entropy changes, and mean ionic activity coefficients were calculated for each solvent mixture and temperature. Relationships of the stand-ard potentials and thermodynamic functons with respect to solvent compositions in the two mixed-solvent systems and the pure solvents were discussed. 

Saboungi, Marie Louise Marr, Jane Blander, Milton. "Thermodynamic Properties of a Quasi-ionic Alloy From Electromotive Force Measurements The Li-Pb System," J. of Chem. Phys. 1978. 

The differences in the hydration of a solnte in H2O and D2O have been extensively stndied by measnring their thermodynamic properties, the change of free energy (AG°t), enthalpy (A//°t), and entropy (AY°t) at the transfer of 1 mol of solnte from a highly dilute solution in H2O to the same concentration in D2O under reversible conditions (mostly 25 °C and atmospheric pressure). Greyson measured the electromotive force (emf) of electrochemical cells of several alkali halides containing heavy and normal water solutions. The cell potentials had been combined with available heat of solution data to determine the entropy of transfer of the salts between the isotopic solvents. The thermodynamic properties for the transfer from H2O to D2O and the solubilities of alkali halides at 25° in H2O and D2O are shown in Table 4. 

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